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Chemical Bonding Explanation

Updated on July 14, 2017

Chemical Bond and Typres of Chemical Bond

Chemical Bond:
The force that binds atom together in a molecule or crystal is called a chemical bond.
Generally, there are two major types of chemical bonds.
Ionic or electrovalent bond
Covalent bond
Ionic or Electrovalent Bond:
A bond that is formed by the complete transfer of electrons OR the attractive force that binds oppositely charged ions together Is known as ionic or electrovalent bond.
Examples:
Formation of NaCl:
In this combination, Na atom transfer one outer most shell electron and become Na+1 ion and Cl atom gains that electron to complete its octet and becomes Cl-1 ion. The attraction that binds Na+ and Cl- ions together is called ionic bond and compound NaCl is called ionic compound.
Na —→ Na+ + e-
2,8,1 2,8
Cl + e- —→ Cl-
2,8,7 2,8,8
Na+ + Cl- —→ Na+Cl-
Formation of MgO:
In this combination, Mg atom loses 2 electrons from its outer most shell to become Mg+2 ion and O atom gain those electrons to complete its octet and becomes O-2 ion. These two positively charged ions form ionic bond and the compound MgO is called ionic compound.
Mg —→ Mg+2 + 2e-
2,8,2 2,8
O + 2e- —→ O-2
2,6 2,8
Mg+2 + O-2 —→ Mg+2O-2
Characteristics of Ionic Compounds:

Ionic compounds are solid at room temperature
They have high melting and boiling point
They don’t conduct electricity in solid state but they conduct electricity in molten form or in aqueous solution
They are usually soluble in polar solvents and insoluble in nonpolar solvents.
Covalent Bond:
The bond that is formed by the mutual sharing of electrons is called covalent bond OR the shared pair of electrons which links the atoms in a molecule is known as covalent bond.
In covalent bond 1 shared pair of electrons Is represented by one short line (—).
Example:
Formation of molecular Chlorine (Cl2):-
The valence shell of chlorine Cl contains seven electrons, and need only 1 electron to complete its octet. So, in the formation of chlorine molecules (Cl2) both atoms of chlorine share 1 electron as shown below



Formation of Molecular Chlorine

Source

Types of Covalent Bond w.r.t No. of Shared Electrons:
Single Covalent Bond:
A covalent bond that is formed by mutual sharing of 1 electron is called single covalent bond. It is represented by a single short line(—).

Double Covalent Bond:
A covalent bond that is formed by a mutual sharing of two electrons is called double covalent bond. It is represented by two short lines(=).

Triple Covalent Bond:
A covalent bond that is formed by the mutual sharing of three electrons is called triple covalent bond. It is represented by three short lines(≡).

Examples for Single, Double and Triple Covalent Bond

Source

Characteristics of Covalent Bond:
They are usually made up of molecules with weak intermolecular forces
They are often gases, liquids or soft solids
They have low melting and boiling point
They donot conduct electricity
They are usually insoluble in polar solvents but soluble in organic solvents
Electronegativity(E.N):
The power of an atom to attract the shared pair of electrons towards itself is known as electronegativity.
For example in HCl molecule, the attraction of Cl atom for electron pair is more than that of H atom. So, Cl is more electronegative than H.
Pauling E.N Table:-
Linus pauling, devised a method to represent E.N values of different elements. He gave Flourine a standard value of electronegativity as 4.0. and the E.N value of other elements are compared with that of F.
Important to Note:
E.N values are expressed by numbers. There is no unit for it.
Non metals have higher E.N values, than metals
Flourine has the highest E.N value 4.0
Cesium has the lowest E.N value 0.7
Ionic Character in a Covalent Bond:-
Ionic character depends upon the difference of E.N values of the bounded atoms. On the basis of it there are three types of bond:
Non-polar Covalent Bond:
If the difference in the E.N values of the bounded atoms in zero, then the bond is pure or non-polar covalent bond.
For Example:
H—H, O=O, N≡N , H3C—CH3 , H2C=CH2 , Cl—N—Cl

Cl
Polar Covalent Bond:
If the difference in the E.N values the bonded atoms is upto 1.7 , then the bond is polar covalent or partially ionic in character.
For Example:
H+8—Cl-8 , H+8—N-8—H+8 , H+8—O-8—H+8

H+8
Electrovalent Bond:-
If the difference in the E.N values of the bounded atom is more than 1.7 than that bond is purely ionic or electrovalent bond.
For Example:
Na+Cl-

Electronnegativity Table

Compound
E.N value of 1st bounded atom
E.N value of 2nd bounded atom
Difference of E.N values
Cl2
Cl (3.0)
Cl (3.0)
0
HCl
H (2.1)
Cl (3.0)
0.9
NaCl
Na (0.9)
Cl (3.0)
2.1

Co-ordinate Covalent Bond or Dative Bond:-
It is a speciall type of covalent bond in which both electrons forming a bond is provided by one atom only.
It is represented by an arrow (→)
The atom which supplies pair of electrons is known as Donor.
The atom which accepts pair of electrons, is known as Acceptor
The pair of electrons transferred is called lone pair of electrons.
Examples:
Formation of Ammonium Chloride:
The N atom of ammonia (NH3) donate electron pair and H+ ion of Hydrogen Chloride(HCl) accept to form NH4Cl.
H N + H+8 Cl

H +
I
H — N → H + Cl- i.e. NH4+Cl-
I
H

Co-ordinate covalent bond
Formation of Hydronium Ion:-
Oxygen O atom of water acts as a donor and hydrogen ion(H+) as acceptor in the formation of H3O+ ion.

H O H+ ——→ H — O → H +
or H3O+
I
H

Metallic Bonding

Metallic Bonding:-
Since metal atoms have less than four valence electrons. These electrons are not confined to any particular atom, but they move freely from one atom to another. So, the atoms must have positive charge. Therefore metals is defined as;
“A substance consisting of positively charged ions, fixed in a crystal lattice with negatively charged electrons moving freely through the crystals.”
These free moving electrons act as cohesive force which hold the atoms together and form metallic bond. So, metallic bond can be defined as;
“It is a combination of electrostatic forces between electron and positive nuclei of atoms.”
Most o the observed properties of metals are because of metallic bonding like ductility, malleability, conductance, etc.
Metallic Bonding in Sodium (Na):
Na metal has one valence electron per atom. It crystallizes in a body centered cubic structure in which each atom is surrounded by eight nearest neighbours. The valence electrons are free to move throughout the crystals resulting in weak metallic bonding. That is why, metals like Na and K have relatively low melting points.
Metallic Bonding in Iron (Fe) and Copper (Cu):
These metals have incomplete valence shells therefore the atoms become covalently bounded to each other throughout the crystals. That’s why iron (Fe) and copper (Cu) are hard and have high melting points.
Explanation of the properties of Metals:-
Conductivity:-
Since electrons are free to move so metals are good conductor of electricity.
When metals are heated, the mobile electron absorb heat and transfer it to neighbouring electrons, this means that metals are good conductor of heat.
Luster:-
The mobile electron readily absorbs light falling upon than and moves to lighter energy level. When they return to their original position. They emit radiations. This is the cause of luster.
Malleability and Ductility:-
The delocalized electrons allow the metal ions to slide over each other. This makes metals malleable which means that the are soft, easily bent and shaped, and can be presses or beaten into thin sheets. Metals are also ductile which means that they can be drawn down into wires
Forces Present Among Molecules:-
There are two types of forces present among the molecules.
Intra-molecular Force:-
A force which held atoms together in a molecule is called intra-molecular force. E.g. ionic and covalent bond.
Intermolecular forces:-
The forces that holds neutral molecules together at certain temperature. These are also called Vander Waal’s forces.
There are three types of inter molecular forces:
Dispersion forces:-
These are weak attractive forces between temporarily polarized atoms (or molecules) due to the varying positive of electrons during their motion around nucleus.
These are also known as London forces because they were first identified by Fritz London. These have small energies i.e. 1-10 kj/mole.
For Example:
When two Ne atoms come extremely close to each other. The electron cloud repel each other and as a result, both atom polarized and dispersion forces are developed.

Dispersion forces Example

Source

Dipole-Dipole Forces

Dipole- Dipole forces:-
A force that generates due to the intraction of the positive end of one molecule with the negative end of the other.
For Example:-
In HCl molecule, Hydrogen is partially positive while chlorine is partially negative i.e. H+6 — Cl-6 . so, dipole- dipole forces acts between the positive end of one Hcl molecule and the negative end of other as shown in figure.

Example

Source

Hydrogen Bonding:-
A dipole-dipole attractive force that exists between two polar molecules containing a hydrogen atom covalently bounded with F, O or N.
A hydrogen bond can have 5% to 10% of the strength of covalent bond. It is represented by dotted line(-----). Water is the best example of hydrogen bonding.

Example

Source

Which bond is more stronger?

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